Chemistry 110
Dr. A. J. Pribula
Predicting the Products of Chemical Reactions
Introduction
Predicting the product(s) which will form in a given set of circumstances (or making a prediction of “no
reaction”) is a topic which often causes difficulty for students in introductory Chemistry courses. Many try to
memorize reaction after reaction, which leads to confusion and frustration. Since over 50 million chemical
compounds are currently known, trying to memorize even a tiny fraction of their possible reactions is a
virtually impossible task—for anyone, not just introductory Chemistry students. A systematic approach to the
task is therefore called for. There are only a small number of fundamental classes of reactions that you need to
be familiar with for this course, and a few common categories of reaction within these classes. Once you realize this, it can go a long way towards organizing and systematizing your approach to this topic. If you train
yourself to recognize certain general reaction patterns, and to classify the reactant(s) in a potential reaction into
classes such as element (metal or nonmetal), compound (ionic or molecular), strong or weak electrolyte, oxidizing agent or reducing agent, acid, base, etc., you have a framework for thinking about this, and the task of predicting the product(s) becomes easier. (You will probably find it useful to review the meanings of these terms
in your textbook.)
The most important classes of reactions encountered in an Introductory Chemistry course are the following:
1) Combination Reactions: Two (or occasionally more) materials react to form a single product. The general
pattern is A + B → C. There are always two or more substances on the reactant side of the equation and only
one on the product side. Some examples are given in sections (1) – (3) below.
2) Decomposition Reactions: The reverse of a combination reaction; one material reacts to form two or more
products. These often need energy in the form of heat to proceed at a noticeable rate. The general pattern is A
→ B + C. There are always two or more substances on the product side of the equation and only one on the
reactant side. Some examples are given in sections (1), (2), (4), and (5) below.
3) Replacement Reactions (also called substitution, exchange, or displacement): One atom or group replaces
or substitutes for another in a compound. There are two sub-classes. In a single replacement reaction, an element replaces the corresponding element in a compound. The general pattern is A + BC → AC + B, in which
the element A is replacing the element B. Some examples are given in section (10) below. In a double replacement (also called a metathesis reaction), two compounds react, and they “trade partners,” often with the
cations and anions trading places with their counterparts in the other compound. The general pattern is AB +
CD → AD + CB. In both of these sub-classes, the number of substances on the reactant side of the equation is
the same as that on the product side.
Two important types of double replacement reactions are found. In an acid-base reaction, a proton (H+ ion)
from the acid is transferred or donated to the base. Often, the base produces OH¯ ions, which combine with
the H+ to form water. In a precipitation reaction, an insoluble ionic compound is formed from soluble reactants
when the cation from one compound combines with the anion from another. Examples of acid-base reactions
are given in sections (6), (7), and (9) below, while examples of precipitation reactions are given in section (8).
4) Oxidation-Reduction (Redox) Reactions: A transfer of electrons (at least a partial one) occurs between the
two reactants. A loss of electrons (increase in oxidation number) is oxidation, while a loss of electrons (or
decrease in oxidation number) is reduction. Many combination and decomposition reactions involve oxidation
and reduction, as do all single replacement reactions. Examples are given in sections (1), (4), (10), and (11)
below.
Some Specific Categories
1. When two elements react, a combination reaction can occur (think: could any other type of reaction occur?), producing a binary compound (that is, one consisting of only two types of atoms). These reactions can
also be described as oxidation-reduction (redox) reactions. (See (11) below.) If a metal and a nonmetal react,
the product is ionic with a formula determined by the charges on the ions the elements form. If two nonmetals
react, the product is a molecular material with polar covalent bonds, with a formula consistent with the typical
valences of the atoms involved. Some pairs of elements may react only slowly and require heating for significant reaction to occur. The chemical reactivity of metals in this type of reaction is similar to the order given in
section (10) below, with the least reactive metals perhaps not reacting at all. Compounds of metals of low
reactivity can decompose when heated, and most binary compounds can be decomposed by the passage of an
electrical current through them (usually in a molten state).
Examples:
K(s) + S8(s) → K2S(s)
(ionic)
Ca(s) + O2(g) → CaO(s)
(ionic)
Pt(s) + O2(g) → no reaction
Al(s) + I2(s) → AlI3(s)
(ionic)
H2(g) + O2(g) → H2O (l)
(covalent)

C(s) + O2(g) → CO(g) or CO2(g)
(covalent)
Δ
I2(s) + Cl2(g) → ICl(l), ICl3(l), or ICl5(s)
(covalent)
(The exact product in the above two cases depends on the relative amounts of reactants, and other conditions.)
HgO(s) → Hg(l) + O2(g)
Δ
NaCl(l) → Na(s) + Cl2(g)
elect.
(NOTE: The above reactions are not balanced, nor were they intended to be. They, like the others in this
handout, are meant only to show the correct formulas for the reactants and products. You may wish to balance
the reactions in the handout as an exercise. After balancing, converting the reactions which occur in aqueous
solution to net ionic form would also be a useful exercise.)
2. Reaction of a metal oxide with water produces a metal hydroxide; that is, a strong base. Reaction of a nonmetal oxide with water produces an oxyacid in which the nonmetal is in the same oxidation state as in the
oxide you started with. All acids and some metal hydroxides are soluble in water, so the result is an aqueous
solution of the acid or base if an excess of water is used (as assumed in the reactions below). Both of these are
combination reactions, and both can be reversed by heating the products. Metal hydroxides decompose on
heating to give the metal oxide and water, and oxyacids decompose on heating to give water and the nonmetal
oxide in the appropriate oxidation state.
Examples:
Na2O(s) + H2O(l) → NaOH(aq)
MgO(s) + H2O(aq) → Mg(OH)2(aq)
SO2(g) + H2O(l) → H2SO3(aq)
P2O5(s) + H2O(l) → H3PO4(aq)
Cl2O5(s) + H2O(l) → HClO3(aq)
HNO3(l) → N2O5(s) + H2O(g)

Fe(OH)3(s) → Fe2O3(s) + H2O(g)

3. Reaction of a metal oxide with a nonmetal oxide gives an oxysalt; reaction of a metal hydroxide with a
nonmetal oxide produces a “hydrogen” oxysalt. This is essentially a reaction of the O2¯ or OH¯ in the metal
compound with the molecular nonmetal oxide. This combination reaction occurs only if no water is present; in
the presence of water, the nonmetal and metal oxides react with the water to produce acid and hydroxide,
respectively (as shown in (2) above), then these react as in (6) below. Typically, this type of reaction does not
involve oxidation or reduction.
Examples:
CaO(s) + SO3(g) → CaSO4(s)
NaOH(s) + CO2(g) → NaHCO3(s)
4. Heating an oxysalt results in decomposition into a metal oxide plus a nonmetal oxide or a metal salt plus
oxygen, or some combination of these two reactions. Some decomposition reactions involve oxidation and
reduction (such as the second and third examples below), but some do not (such as the first example).
Examples:
CaCO3(s) → CaO(s) + CO2(g)

KClO3(s) → KCl(s) + O2(g)

Pb(NO3)2(s) → PbO(s) + NO(g) + NO2(g) + O2(g)

5. Heating a hydrated material initially causes a decomposition reaction to produce the anhydrous compound
and water. Further heating may yield further decomposition, depending on the material. (See (2) and (4)
above.) Most binary compounds are stable to heat and will not decompose further. (See (1) above.)
Examples:
CaCl2·6H2O(s) → H2O(g) + CaCl2(s);
followed by

CaCl2(s) → no reaction

H2C2O4·2H2O(s) → H2O(g) + H2C2O4(s);
followed by

H2C2O4(s) → H2O(g) + CO(g) + CO2(g)
Δ
CuSO4·5H2O(s) → H2O(g) + CuSO4(s);
followed by

CuSO4(s) → CuO(s) + SO3(g)
(requires strong heating)

6. Reaction of an acid with a base gives a salt plus water. The cation in the salt comes from the base; the anion comes from the acid. (You can remember this if you keep the consonants (c, b) together and the vowels (a,
a) together.) The base may be a metal hydroxide, a metal oxide, or a weak base such as NH3. (Note that weak
bases are a special case, in that they do not yield water as a product.) The acid and/or base may be pure solids,
liquids, or gases, or in aqueous solution. The oxidation states of the anion of the acid and cation of the base
normally remain unchanged. This is a special case of a double replacement reaction.
Examples:
HCl(aq) + Ca(OH)2(aq) → CaCl2(aq) + H2O(l)
H2SO4(aq) + Fe(OH)3(s) → Fe2(SO4)3(aq) + H2O(l)
NH3(g) + HC2H3O2(l) → NH4C2H3O2(s)
Al2O3(s) + HClO4(aq) → Al(ClO4)3(aq) + H2O(l)
7. Ammonium salts react with metal hydroxides and oxides in an acid-base reaction to produce ammonia.
This is essentially the reverse of one of the reaction types mentioned in (6) above. Either or both of the reactants may be a pure material or in aqueous solution. Salts of other weak bases react in a similar fashion.
Examples:
NH4Cl(aq) + KOH(aq) → NH3(g) + H2O(l) + KCl(aq)
NH4NO3(aq) + CaO(s) → NH3(g) + H2O(l) + Ca(NO3)2(aq)
8. Reaction of solutions of two soluble strong electrolytes (acids, bases, or salts) with one another can give a
precipitate of an insoluble ionic compound formed by a double replacement reaction (also called a metathesis).
Whether or not a precipitate forms depends on the exact combination of materials used. To make a prediction
as to whether a reaction will take place or not, you must know the solubility rules for common salts (Chang,
5/e, p. 98 (you should add acetates to the list of soluble compounds in the group with nitrates, etc.); GenChem
lab manual, Appendix 11, p. A-51; Dennison, Topping, and Caret, 5/e, p. 130). Some combinations of compounds may give oxidation-reduction reactions (see (11) below), but most do not.
Examples:
CaCl2(aq) + K2CO3(aq) → CaCO3(s) + KCl(aq)
AgNO3(aq) + FeCl3(aq) → AgCl(s) + Fe(NO3)3(aq)
H2SO4(aq) + BaBr2(aq)  BaSO4(s) + 2 HBr(aq)
Cr2(SO4)3(aq) + KOH(aq)  Cr(OH)3(s) + K2SO4(aq)
but:
NiSO4(aq) + MgI2(aq) → no reaction
(NiI2 and MgSO4 are both soluble strong electrolytes)
Al(NO3)3(aq) + Pb(C2H3O2)2(aq) → no reaction
(Al(C2H3O2)3 and Pb(NO3)2 are both soluble strong electrolytes)
HCl(aq) + ZnBr2(aq)  no reaction
(HBr and ZnCl2 are both soluble strong electrolytes)
(Changing the above into net ionic form (after balancing them, of course) would be a good way to show that
the last three indeed result in no reaction.)
9. Reaction of the salt of a weak acid (that is, a compound containing the anion of a weak acid) with a strong
acid produces the weak acid and a salt. This is another example of an acid-base (and double replacement) reaction, in addition to the ones given in (6) and (7) above. The original salt of the weak acid may be either a pure
solid or in aqueous solution. The cation in the salt formed as the product comes from the weak acid salt; the
anion in the product salt comes from the strong acid. In many cases, the weak acid produced is unstable and
decomposes to give the oxide of a nonmetal and water (see (2) above). This is especially true if the nonmetal
oxide is a compound of limited solubility in water such as SO2, CO2, or the nitrogen oxides. The best-known
examples of this type of reaction involve carbonates, hydrogen carbonates, sulfides, phosphates, acetates, and
sulfites, but many other examples are known as well. Normally, these reactions do not involve oxidation or reduction. Production of bubbles when acid is added to a sample is a commonly-used “spot test” for carbonate or
hydrogen carbonate ion.
Examples:
BaCO3(s) + HBr(aq) → BaBr2(aq) + H2O(l) + CO2(g)
NaHCO3(aq) + H2SO4(aq) → Na2SO4(aq) + CO2(g) + H2O(l)
MgS(s) + HCl(aq) → H2S(g) + MgCl2(aq)
Ca3(PO4)2(s) + HCl(aq) → CaCl2(aq) + H3PO4(aq)
Zn(C2H3O2)2(aq) + HBr(aq) → ZnBr2(aq) + HC2H3O2(aq)
K2SO3(aq) + HNO3(aq) → KNO3(aq) + SO2(g) + H2O(l)
10. Reaction of an element with a compound often gives a single replacement reaction in which a nonmetallic
element can replace a combined nonmetal, and a metallic element can replace a combined metal, or hydrogen
from an acid. As a general rule, a more active (reactive) element will replace a less active (reactive) element
from its compounds. In general (but with many exceptions), the most reactive nonmetals are found to the upper right in the periodic table, and the most reactive metals are found to the left. The order of reactivity of the
halogens is F2>Cl2>Br2> I2. (Other nonmetals do not participate in this type of reaction.) For hydrogen and the
more common metals, the order of reactivity (the activity series) is
Li>K>Ca>Na>Mg>Al>Zn>Cr>Fe>Ni>Sn>Pb>H2>Cu>Hg>Ag>Pt>Au
(This is also discussed in Chang, 5/e, p. 114, and in the GenChem lab manual, Appendix 12, p. A-53.) In
these two series, one element can replace another one to its right in the series. Metals to the left of H2 in the
series can replace H+ from acids. The very reactive metals (Li, K, Ca, Na) can replace H+ from cold water;
metals of intermediate reactivity (Mg, Al) can replace H+ from hot water or steam. All single replacement
reactions can also be categorized as oxidation-reduction (redox) reactions (see (11) below).
Examples:
Al(s) + NiSO4(aq) → Al2(SO4)3(aq) + Ni(s)
Fe(s) + HBr(aq) → FeBr3(aq) + H2(g)
Cl2(g) + KI(aq) → KCl(aq) + I2(s)
Na(s) + H2O(l) → NaOH(aq) + H2(g)
Zn(s) + Cu(NO3)2(aq) → Cu(s) + Zn(NO3)2(aq)
but:
Ag(s) + HClO4(aq) → no reaction
(Ag is less reactive than H2.)
Br2(l) + ZnCl2(aq) → no reaction
(Br2 is less reactive than Cl2.)
Sn(s) + H2O(l) → no reaction
(Sn is not reactive enough to replace H+ from H2O.)
Pb(s) + CrF3(aq) → no reaction
(Pb is less reactive than Cr.)
11. Compounds containing atoms in high oxidation states often act as oxidizing agents; compounds containing
atoms in low oxidation states often act as reducing agents. For most elements, the (old) group number of the
atom in the periodic table (the second digit of the “new” group number) gives the highest oxidation state
possible for that element. (However, note that some elements can only be oxidized to their highest possible
oxidation state with great difficulty, and some not at all.) For nonmetals, the lowest oxidation state possible is
given by the (old) group number minus eight (or the “new” group number minus eighteen). Elemental metals
most often act as reducing agents (they are oxidized); nonmetals frequently act as oxidizing agents (they are
reduced).
For the representative elements (also called the main-group elements; those in Groups 1, 2, 13-18 in the periodic table), oxidation states most often are two units apart. For example, Sn forms Sn(II) and Sn(IV); Br forms
Br¯, Br(I), Br(III), Br(V), and Br(VII). For the transition elements, (i.e., those in Groups 3-12 of the periodic
table), oxidation states are often one unit apart, but can be in almost any relationship to one another. For the
transition metals (as well as some of the main-group metals), the common oxidation states (charges on their
ions) must be memorized. For example, Fe forms Fe2+ and Fe3+; Cu forms Cu+ and Cu2+, Pb forms Pb2+ and
Pb4+, etc. Some of the transition elements form oxyanions as well as cations. For example, Mn forms Mn2+,
Mn3+, MnO42¯, and MnO4¯; Cr forms Cr2+, Cr3+, CrO42¯, and Cr2O72¯.
Any atom in its highest possible oxidation state can only act as an oxidizing agent; any atom in its lowest
possible oxidation state can only act as a reducing agent. Atoms in intermediate oxidation states can be either
oxidized or reduced; that is, they can act as either reducing or oxidizing agents. Some of the oxidizing agents
most commonly encountered are H2O2, MnO4¯, CrO42¯, Cr2O72¯, HNO3, and the halogens. Some of the more
common reducing agents are elemental H2, metals, carbon, and I¯.
In predicting products of oxidation-reduction reactions, don't forget their name—oxidation and reduction must
occur simultaneously! It is impossible for oxidation to occur without reduction or vice versa.
Examples:
Sn2+(aq) + F2(g) → Sn4+(aq) + F¯(aq)
Mn2+(aq) + BiO3¯(aq) → Bi3+(aq) + MnO4¯(aq)
(note that the Bi is in its highest possible oxidation state in BiO3¯)
K(s) + P4O10(s) → K3PO3(s)
(note that P is reduced from P(V) to P(III))
MnO4¯(aq) + I¯(aq) → Mn2+(aq) + I2(aq)
(note that Mn is in its highest oxidation state in MnO4¯, and that I¯ is in its lowest)
CuS(s) + HNO3(aq) → Cu(NO3)2(aq) + S8(s) + NO2(g)
(note S2¯ → S0 and N(V) → N(IV))
Fe2O3(s) + C(s) → CO2(g) + Fe(s)
Δ
(Other examples of redox reactions can be found in sections (1), (4), and (10) above.)
Reaction 110.doc
9/09